Atomic Structure Notes - Boys Ranch HS...
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Atomic Structure Notes Notes Packet for Chem Website – Unit 2: Atomic Structure (Final update)
Transcript of Atomic Structure Notes - Boys Ranch HS...
Atomic Structure Notes
Notes Packet for Chem Website – Unit 2: Atomic Structure (Final update)
Particle Theory
John Dalton
• John Dalton
• 1808
• England
• Described atoms as
tiny particles that
could not be divided.
Thought each
element was made of
its own kind of atom.
• J. J. Thomson
• 1897
• England
• Thompson discovered
that electrons were
smaller particles of an
atom and that they
are negatively
charged.
• Cathode Ray
Experiment
Discovery of Electrons
• Ernest Rutherford
• 1911- England
• Isolated the positive particles
in an atom. Decided that the
atoms were mostly empty
space, but had a dense core.
• Gold Foil Experiment
Atomic Structure I
Ernest
Rutherford
Niels Bohr
1913
England
Proposed that electrons traveled in fixed paths around the nucleus. Scientists still use the Bohr model to show the number of electrons in each orbit around the nucleus.
Atomic Structure II
Atoms
An atom is the smallest complete part of an element that maintains the properties of the element.
Combine to form molecules
They make up everything we see, hear, touch, smell and feel.
They are so small, one cell from the body contains 100 trillion atoms.
Parts of an Atom
the atom is made of three subatomic particles:
- protons (positive charge), p+
- neutrons (no charge), nO
- electrons (negative charge), e-
Anatomy of an Atom
Nucleus
protons: mass = 1amu
neutrons: mass = 1amu
~1/1000 of the atom’s diameter
Cloud
electrons: mass ~ 1/1000 amu ~ 0
does not factor into atomic mass
Electron Cloud
Protons (+) and Neutrons (o) inside
Nucleus
Electrons (-) outside
Reading the Periodic Table
Atomic Number = number of protons
- This defines what the element is!!!
Atomic Mass = protons + neutrons
- Remember, electron mass ~0amu
- round to the nearest whole number
QuickTime™ and aTIFF (Uncompressed) decompressor
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are needed to see this picture.
Ions & Isotopes
Since it is the number of protons, or the atomic number, that
defines what the element is….
What happens…
If p = e? Neutral atom
If p < e? Ion with a negative charge (anion)
If e < p? Ion with a net positive charge (cation)
If extra neutrons? Heavier isotopes, possibly unstable!
The Role of Electrons
How are electrons arranged in
atoms?
• Usually, number of e- equals number of
p+ (the atom is neutral)
Energy Levels
• e- are in the electron
cloud, which is divided
into energy levels
• The first energy level is
the closest to the
nucleus
Energy levels (contd.)
• 2n2 = the number of electrons that will
fit in an energy level (given that n = the energy level number)
– 1st = 2 e-
– 2nd = 8 e-
– 3rd = 18 e-
– 4th = 32 e-
Valence Electrons
• Valence electrons are
the electrons found in
the outermost energy
level of the atom
• Maximum number of
valence e- is 8 (with
exception of first
energy level)
Bohr Model
Write number of protons and neutrons in the middle
Fill in the energy levels (shown below) with however
many electrons the given atom has.
What element is this?
Is it neutral?
Using the periodic table
to count valence e-
Group # 1 2 13 14 15 16
17 18
• Groups (columns) are numbered 1-18,
• number of valence e- is the last number
(for example: group 18 would have 8, group 2 would have 2)
Octet
• When atom has eight
valence electrons
• Octets are chemically
stable (happy)
• Atoms want to be
stable, so they will
gain, lose or share
electrons to get an
octet
Ions
• Ions are atoms that have gained or lost
an electron to get an octet
– Gain an e- = get more negative
– Lose an e- = get more positive
(BREAK HERE TO DO THE IONS
LAB WITH THE CPO BOARDS)
Lewis Dot Symbols
Practice with Valence Electrons
(Teacher note: to be completed as time
permits)
Valence Electrons
• Valence electrons are
the electrons found in
the outermost energy
level of the atom
• Maximum number of
valence e- is 8 (with
exception of first
energy level)
Using the periodic table
to count valence e-
- Groups
are
numbered
1-18
- number of
valence e-
is the last
number
Group # 1 2 13 14 15 16
17 18
Octet
• When atom has eight
valence electrons
• Octets are chemically
stable (happy)
• Atoms want to be
stable, so they will
gain, lose or share
electrons to get an
octet
To Draw a Lewis Dot Symbol:
1. Write the symbol for the atom
2. Find the number of valence electrons (use Periodic Table)
3. For every valence electron, draw dot around the symbol
Example #1: Sodium
Na *Sodium has 1 valence
electron, so we draw one dot.
Arrangement of Electrons
• Pretend there’s a box around the symbol
• Draw the first e- on one side of the box, then rotate to the next side and draw another
• Keep rotating until you’ve drawn them all
Example #2: Carbon
C *Carbon has 4 valence
electrons
Arrangement of Electrons
• Up to two electrons can be on each “side”
• Valence e- prefer to be in pairs
• (one of the reasons atoms bond with other atoms is to pair up their valence e-)
Example #3: Sulfur
S *Sulfur has 6 valence
electrons
Can use drawing to
determine the charge!
Sulfur will gain 2 e- to
get to 8, so charge is
-2.
Isotopes
• an isotope is an element with a
different
number of neutrons (and a different
mass)
• neutrons = atomic mass – atomic #
Your turn…
• Do numbers 1-3 on your Isotopes
Practice Problems.
• I will come by and stamp your work.
You have 8 minutes.
Why isn’t atomic mass a whole number?
• The atomic mass of an element is a
weighted average calculated using the
mass and relative abundance (a.k.a.
%)of each isotopes.
• This is similar to how your GRADES are
calculated (Tests = 30%, Daily work =
50%, Quizzes = 20%)
Average Atomic Mass
• atomic mass is calculated using the
mass of each isotope times its relative
abundance
• ex. Carbon (Atomic mass = 12.01)
– 98.93% of carbons have a mass of 12 amu
– 1.07% of carbon atoms have a mass of 13
amu
Calculating Average Atomic
Mass 1. Convert % abundance to decimal form.
(by moving the decimal two places to the left)
2. Multiply the decimal by its respective atomic
mass
3. Add the products.
Review Questions
• In an isotope, the number of which
subatomic particle changes?
• Calculate the number of protons and
neutrons in Sulfur-35.
Electron Movement
Role of Electrons – Part III
Warm-up 9/19 and 9/20
1. Draw a Bohr model of oxygen.
2. Then draw a lewis dot
structure.
3. What will oxygen do to be
happy and stable?
4. What will its charge be?
Warm Up: 11 Oct. 2013
• Read the instructions for the flame test
lab.
• Write down any safety considerations or
rules you can think of that would keep
us safe during this lab.
Flame Test Lab
• NOTES:
– Be very careful when working with bunsen
burners in the lab! Follow safety rules.
– Keep the samples clean. Use new Q-tips and
avoid touching them to the burner. If
contaminated, you will not see the colors.
– Do not touch the salt directly. Wear gloves or
wash your hands after touching any lab
equipment.
Electron location
• e- prefer to be in the lowest energy level available, called the “ground state”
• If energy is added, e- can move up to a higher energy level, called the “excited state”
Energy transfer
• Energy must be absorbed or released by the atom in order to move e- between levels
e- move down when e- move up when
energy is released energy is added
Photons • Photon – a discrete bundle of energy
• We see the energy of photons as colored
light based on their wavelength and
frequency
Light Spectrum
• The energy difference between the atom’s energy levels determine the color of the light
• Higher energy is closer to the purple end of the spectrum
Wavelength
Frequency
ENERGY
Bright Line Spectra
• Released energy
shows up as a
colored line
• Absorbed energy
shows up as a black
line
Atomic fingerprints
• Each element has a distinct set of photons that can be absorbed or released
• Astronomers use these spectral lines to identify which elements are present in stars
Flame Test Lab
• We can see the most abundant photons
by doing a flame test
• Metal salts burn a distinctive color
The Periodic Table
1IA
18VIIIA
11H
1.00797
2IIA
Periodic Table 13IIIA
14IVA
15VA
16VIA
17VIIA
2He4.0026
23
Li6.939
4Be9.0122
5B10.811
6C
12.0112
7N
14.0067
8O
15.9994
9F
18.9984
10Ne20.179
311
Na22.9898
12Mg24.305
3IIIB
4IVB
5VB
6VIB
7VIIB
8 9VIIIB
10 11IB
12IIB
13Al26.9815
14Si28.086
15P
30.9738
16S32.064
17Cl35.453
18Ar39.948
419K39.102
20Ca40.08
21Sc44.956
22Ti47.90
23V50.942
24Cr51.996
25Mn54.9380
26Fe55.847
27Co58.9332
28Ni58.71
29Cu63.54
30Zn65.37
31Ga65.37
32Ge72.59
33As74.9216
34Se78.96
35Br79.909
36Kr83.80
537
Rb85.47
38Sr87.62
39Y88.905
40Zr91.22
41Nb92.906
42Mo95.94
43Tc[99]
44Ru101.07
45Rh102.905
46Pd106.4
47Ag107.870
48Cd112.40
49In114.82
50Sn118.69
51Sb121.75
52Te127.60
53I
126.904
54Xe131.30
655Cs132.905
56Ba137.34
57La138.91
72Hf178.49
73Ta180.948
74W183.85
75Re186.2
76Os190.2
77Ir192.2
78Pt195.09
79Au196.967
80Hg200.59
81Tl204.37
82Pb207.19
83Bi208.980
84Po[210]
85At[210]
86Rn[222]
787Fr[223]
88Ra[226]
89Ac[227]
104Ku[260]
105 106 107 108 109
Origin of the Periodic Table
• originally organized
by Dmitri
Mendeleev in the
1850s
• arranged the known
elements according
to atomic mass
Origin of the Periodic Table • Mendeleev noticed that when you put the
elements in rows of 8 that all the columns
seemed to have similar elements in them
Origin of the Periodic Table
• Henri Moseley
discovered that
each element has a
unique atomic
number
• current periodic
table is organized
by atomic number
Organization of the Periodic
Table
1) By increasing atomic number
2) Divided into metals, nonmetals, and
metalloids
3) Divided into “A” and “B” type elements
4) Into rows and columns
Metals vs. Nonmetals
• Metals are on left of “staircase” line
• Nonmetals are on right of “staircase”
line
• Metalloids (properties of both metals
and nonmetals) touch the “staircase”
line
What do the As and Bs mean?
• “A” group = representative elements
• “B” group = transition metals
Representative vs. Transition
• For the “A” elements you can tell the
valence electrons by looking at the
Roman numeral
• The “B” elements are called transition
elements because they can have many
valences, (trans = change) -- we can’t
use the periodic table to predict!
Columns
• Called groups or families
• Elements have similar chemical
properties
Family: arranged vertically down the periodic
table
(1- 18 or 1-8 A,B)
*have the same number of e- in the outer most or valence
shell 1IA
18VIIIA
12IIA
13IIIA
14IVA
15VA
16VIA
17VIIA
2
33IIIB
4IVB
5VB
6VIB
7VIIB
8 9VIIIB
10 11IB
12IIB
4
5
6
7
Alkali Family:
1 e- in the valence shell
Halogen Family:
7 e- in the valence shell
Rows
• Called periods or series
• The row is the number of e- energy
levels
• Elements have similar masses
Periods: arranged horizontally across periodic
table
(rows 1-7)
*have same number of valence shells 1IA
18VIIIA
12IIA
13IIIA
14IVA
15VA
16VIA
17VIIA
2
33IIIB
4IVB
5VB
6VIB
7VIIB
8 9VIIIB
10 11IB
12IIB
4
5
6
7
2nd Period
6th Period
Review Questions:
1) Who originally organized the elements into what later became the periodic table?
2) What property did he use to organize the periodic table?
3) How are the elements on the periodic table arranged now?
4) How do you find the number of valence electrons for representative elements?
5) How can you tell which elements are metals and which are non-metals?
Unit 4 – Bonding and
Electrons
The Role of Electrons
How are electrons arranged in
atoms?
►Usually there are as many electrons as
protons (the atom is neutral)
►Electrons are found in the electron
cloud
►The electron cloud is divided into
energy levels
Ions
►Ions are atoms that have gained or lost an
electron
►Gain an electron = get more negative
►Lose an electron = get more positive
Energy Levels
►The electron cloud is divided into energy levels
►The farther away an electron is from the nucleus, the higher its energy level
►The first energy level is the closest to the nucleus
Energy levels (contd.)
►The energy levels are like the layers of an
onion, as they get farther from the nucleus
they get larger in diameter
►Each increasing energy level holds an
increasing amount of electrons
2n2 = the number of electrons that will fit in an
energy level
(given that n = the energy level number)
Energy levels
►The first energy level
holds two electrons
►The second energy
level has 8 electrons
►The third energy level
has 18 electrons
►The forth energy level
has 32 electrons
Energy sublevels
►It’s important to know that energy levels
can overlap to form sublevels
►In fact, each energy level is subdivided
into regions called electron orbitals
Valence Electrons
►Valence electrons are
the electrons found in
the outermost energy
level of the atom
►Valence electrons are
responsible for forming
chemical bonds
Valence electrons
►The maximum number
of valence electrons an
atom can have is 8
►Exception: the first
energy level has only 2
electrons
Valence Electrons (contd.)
►You can tell how many valence electrons an
atom has by the Roman numeral above the
group it is in (for example: group IV has 4
valence electrons)
►If the groups are numbered 1-18, then the
number of valence electrons is the last number
of the group number (for example: group 18
would have 8, group 2 would have 2)
Octets
►When an atom has eight valence electrons it has an octet
►Octets are chemically stable
►Atoms want to be stable, so they will gain, lose or share electrons to have an octet (BONDING)
Lewis Dot Symbols
Practice with Valence Electrons
(Teacher note: to be completed as time
permits)
To Draw a Lewis Dot Symbol:
1. Write the symbol
for the atom
2. Find the number of
valence electrons
(use Periodic
Table)
3. For every valence
electron, draw dot
around the symbol
Example #1: Sodium
Na *Sodium has 1 valence
electron, so we draw one
dot.
Arrangement of Electrons
• Pretend there’s a box
around the symbol
• Draw the first e- on
one side of the box,
then rotate to the next
side and draw
another
• Keep rotating until
you’ve drawn them all
Example #2: Carbon
C *Carbon has 4 valence
electrons
Arrangement of Electrons
• Up to two electrons
can be on each
“side”
• Valence e- prefer to
be in pairs
• (one of the reasons
atoms bond with
other atoms is to
pair up their valence
e-)
Example #3: Sulfur
S *Sulfur has 6 valence
electrons
Can use drawing to
determine the charge!
Sulfur will gain 2 e- to
get to 8, so charge is
-2.
Types of Chemical Bonding
Quick Check
Role of Electrons Quick Check
#1
Role of Electrons Quick Check
#2
Which of the following is used to draw a
Lewis structure of either a molecule or
atom?
A bond length between 2 atoms
B electronegativity of an atom
C number of valence electrons
D atomic mass or masses
Role of Electrons Quick Check
#3
Choose the correct
Lewis Dot diagram for
Carbon.
Three Main Types of Bonding
• Ionic bonding –a metal and a nonmetal
• Covalent bonding – a nonmetal and a
nonmetal
• Metallic bonding –a metal and a metal
Cop
y!
What type of bond is it?
Find the following elements on the periodic
table and determine their bond type:
a) Ca & S
b) Li & Pt
c) C & Cl
Bonding and Electronegativity
• Large differences in electronegativity = IONIC bonds
• Small differences in electronegativity = COVALENT
bonds
Ionic Bonding
• Metal and non-metal (one is way more
electronegative than the other)
• Metal loses e- and becomes positive
(“cation”)
• Nonmetal gains e- and becomes
negative (“anion”)
• Opposites attract: the ions bond via
opposite charges
Covalent Bonding
• SHARE electrons (the outer energy
levels overlap)
• 1, 2, or 3 electrons shared (single,
double, or triple bond)
• (We will talk Lewis Structures tomorrow)
Metallic Bonding • Electron sea model – atoms of a metal are
fixed in position, but outer electrons move
freely (due to orbitals)
• Delocalized electrons – electrons that
don’t belong to a single atom or bond
• Delocalized e- are attracted
to nucleus (+ charge) =
metallic bond
• Delocalized electrons
cause metals to be
excellent conductors
• Can move heat and
electric charge through
out a metal quickly
Conductivity
By Jan Harenburg (own fotography) [CC-BY-3.0]
• Metals can be changed in shape without causing the crystal
lattice to break
– Malleability – ability to be formed into sheets
– Ductility – ability to be drawn into wire
Malleability and Ductility
I, Daniel Schwen [GFDL] By kurtsik (Forjaria.) [CC-BY-SA-3.0]
Review Questions: 1) Why do atoms bond?
2) Which type of bond shares electrons?
3) How do we represent shared electrons?
4) What is a cation?
5) What type of bond has mobile electrons?
6) Does the metal or the nonmetal become the anion?
7) What is a triple bond?
8) Why are metals such good conductors?
9) What is ductility?
10)What type of bond is this:
Draw and label • In your notes draw each diagram, label
each as covalent, ionic, or metallic:
Today’s Objective(s):
• Electron Configurations (6E) –
Express the arrangement of electrons in
atoms through electron configurations
and Lewis valence electron dot
structures.
• Dot Structures (7C) – Construct
electron dot formulas to illustrate ionic
and covalent bonds.
To Draw a Lewis Dot Symbol:
1. Write the symbol
for the atom
2. Find the number of
valence electrons
(use Periodic
Table)
3. For every valence
electron, draw dot
around the symbol
Example #1: Sodium
Na *Sodium has 1 valence
electron, so we draw one
dot.
Arrangement of Electrons
• Pretend there’s a box
around the symbol
• Draw the first e- on
one side of the box,
then rotate to the next
side and draw
another
• Keep rotating until
you’ve drawn them all
Example #2: Carbon
C *Carbon has 4 valence
electrons
Arrangement of Electrons
• Up to two electrons
can be on each
“side”
• Valence e- prefer to
be in pairs
• (one of the reasons
atoms bond with
other atoms is to
pair up their valence
e-)
Example #3: Sulfur
S *Sulfur has 6 valence
electrons
Can use drawing to
determine the charge!
Sulfur will gain 2 e- to
get to 8, so charge is
-2.
Part 2: Lewis Dot Structures
• Lewis Dot
Structures are used
to depict basic
structures of
covalent
compounds
Steps to Writing Lewis Dot
Structures
• Step 1: Figure out the
skeletal structure
– the least electronegative
atom goes in the middle (the
“central atom”)
– Hydrogen and halogens will
occupy end positions (only
one bond will go to them)
Example 1: Methane (CH4)
Steps to Writing Lewis Dot
Structures
• Step 2: Total the
number of valence
electrons for all
atoms
Steps to Writing Lewis Dot
Structures
• Step 3: Draw a
single bond
connecting the
atoms.
– For each bond
you draw,
subtract 2
valence
electrons from
your total
Steps to Writing Lewis Dot
Structures
• Step 4: Use the
remaining
electrons to
complete octets.
– Remember,
hydrogen only
needs 2 ve- to
have a full outer
energy level -- a
single bond to H is
enough!
Example 2: Ammonia (NH3)
Steps to Writing Lewis Dot
Structures • Step 5: Check for octets. If every
atom now has an octet. You’re done.
If not, go to step 6.
Octet
s!
Octet
s!
Steps to Writing Lewis Dot
Structures
• Step 6: Use
double or triple
bonds to
complete octets
for any atoms
that don’t have
them.
Example 3: Carbon Dioxide (CO2)
Practice!
• Try each of these,
then compare your
structure with your
lab partner’s.
1. CF4
2. Cl2
3. SO2
4. N2
Warm Up: 29 Oct. 2012
• On your own paper . . .
– Draw the Lewis Dot SYMBOL for carbon
by itself
– Draw a dot structure representing the
IONIC bond in NaCl
– Draw the Lewis Dot Structure for the
covalent compound C2H4
Molecular Geometry
Warm-up (in your notebook):
Draw the Lewis Dot Structure of S2:
• Remember the Steps!
– Draw the skeletal structure
– Count the valence e-
– Draw in single bonds (and subtract 2e- for
each)
– Use any extra e- to finish octets
– Check for octets
– If needed, share e- in double or triple bonds
Review: Lewis Dot Structures
• We used Lewis Dot Structures as a 2-D
model of how atoms would bond to
each other when making molecules
• Just like a sketch or a photograph
doesn’t fully convey what you look like,
Lewis Dot Structures don’t show a 3-D
model of a molecule
What is molecular Geometry? • Molecules have shapes
• The types of bonds and
location of lone pairs
determines the shape of
the molecule.
VSEPR Theory Basics • Electrons repel each other, so pairs of
electrons want to be as far apart as
possible
• Bonds and lone pairs around a central
atom will space themselves as far apart as
possible
• It is repulsion that gives molecules their
shape (VSEPR = valence shell electron
pair repulsion)
• Check it out: Molecular Geometry
Simulation
Mo
lecu
lar
Ge
om
etr
ies
“Domains”
• A domain refers to either a lone pair or
any bonding pair (single, double, or
triple bond)
• Tip: the number of domains around a
central atom is how many “things” are
touching it
If an atom has 4 Domains
with… • 4 bonded atoms TETRAHEDRAL
• 3 bonded atoms TRIGONAL PYRAMIDAL
• 2 bonded atoms BENT
If an atom has 3 Domains
with…
• 3 bonded atoms TRIGONAL
PLANAR
• 2 bonded atoms BENT
If an atom has 2 Domains
with…
• 2 bonded atoms LINEAR
Bonding vs. Non-Bonding
Electrons • Electrons in lone
pairs (non-
bonding
electrons) have a
stronger repelling
force than
bonding electrons
• example:
NH4 vs. NH3
Wedge Drawings:
• Wedge drawings can be used to depict a 3D
arrangement on your paper
Review: Valence electrons
• Valence electrons are the electrons found in the outermost energy level of the atom
• Valence electrons are responsible for forming chemical bonds
• TO THE RIGHT: Atoms of nitrogen have 5 valence electrons
Review - Octets
• When an atom has eight valence electrons it has an octet
• Octets are chemically stable
• Atoms want to be stable, so they will gain, lose or share electrons to have an octet (BONDING)
Why do atoms bond?
• We learned that octets (a full valence
electron shell) are very stable. Atoms
desire stability, so they will combine
with other elements in order to achieve
a full octet of valence electrons.
• To become more stable by filling their
outer energy level with electrons (8)
Valence Electrons and the Periodic
Table
• The periodic table is divided into groups
which are columns of elements with
similar properties
• Elements in the same group will have
the same number of valence electrons
(exception: transition metals)
PROPERTIES OF IONIC COMPOUNDS
1.Crystalline solids – usually hard crystals
2.High melting and boiling points –
because of strong forces in the bond
3.Soluble in water –
water molecules attract the ions
4.Conduct electricity (electrolyte)
COVALENT COMPOUNDS:
1. Soft and brittle
2. Low melting and boiling points –
because covalent bonds are weak
3. Insoluble in water –
except if dipole moment
4. Do NOT conduct electricity -
non-electrolyte
Card Sort: Instructions (in pairs . .
.)
1. Arrange the cards with Bohr models in a way that makes sense to you.
2. Blank cards. Draw the appropriate Bohr models to fill in the blanks.
3. Electron Configurations: take the small slips of paper with codes on them, and figure out a way that they will line up with the Bohr models.
4. Predict what the code should look like for the blank cards. Write the codes on these cards.
5. Write your names on the cards you filled out. Staple them together and turn them in with your analysis questions.
Analysis Questions: on a sheet of paper, answer these questions . . .
1. What do the circles on the Bohr represent?
2. Explain what the code means:
– What does the coefficient stand for?
– What does the letter stand for?
– What does the superscript stand for?
3. What is the maximum number of electrons:
– For the “S” orbital?
– For the “P” orbital?
4. How did your group figure out how to match the electron configurations to the Bohr models?
Electron Configurations Video
• http://education-
portal.com/academy/lesson/electron-
configurations-in-atomic-energy-
levels.html
• Take Notes!
OT: Review Videos!
• Valence Electrons
• Ionization Energy
• Electronegativity – talks about bonding
Electron Configuration
Today’s Objectives
• Be able to write the
electron
configuration for any
atom 1-20
• Example:
• Be able to draw
electron orbital
diagrams for any
atom 1-20
• Example:
Exit Ticket: 16 Oct. 2012 • Randomly pick an element from the beaker.
• Write down what element you picked.
• Draw the electron configuration for your
element
• Extra credit: if you have time, draw the orbital
diagram for your element
Electron Configuration
• The electron configuration is notation
which shows how the electrons are
distributed among atomic orbitals and
energy levels.
• Example:
1s2
Helium
Decoding 1s2
• "1" = the coefficient stands for the energy level
• "s" = us that helium’s electrons are in an spherical orbital (sublevel)
• "2" = the superscript is the total number of electrons in that orbital or sub-shell.
Energy Levels
• The energy level is the same as the row
number
1s2
Sublevels
• The first sublevel = s sublevel.
• The second sublevel = p sublevel.
• The third sublevel is called a d sublevel
and the fourth sublevel is called an f
sublevel.
Sublevels and Orbitals
• An orbital is a space that can be
occupied by up to two electrons.
• Each type of sublevel holds a different
number or orbitals, and therefore, a
different number of electrons
II. Sublevels and Orbitals
Total Orbitals and Electrons
• n2 = formula for how many orbitals
are in an energy level
• 2 n2 = formula for how many
electrons are in an energy level
(total)
Total Orbitals and Electrons
V. Order of Filling
Sublevels with Electrons
Friday, October 19th
• Go ahead and log on to the computer
number that matches your desk number
• While you are trying to log on, go ahead
and start the warm up.
Electron Configuration Review 1. Which orbitals are sphere-shaped?
2. Which orbitals are dumbell-shaped, and
arranged along the x, y, and z axis?
3. Which of the following orbitals doesn’t
exist:
4d 3s 4p 3f
1. Which group on the periodic table has d-
block electron configurations that vary?
5. Which of the 2 outermost orbitals MUST BE
FILLED to satisfy the octet rule?
6. How many TOTAL electrons can fit on the third
energy level?
7.
Electron Configuration Review
Lewis Dot Structures
Modeling Covalent Compounds
Metallic Bonds Reading
• Read pages
427-429
• Do Problems
15-18 on
page 429
Warm Up: 1 Nov. 2012
• Get out your review
from yesterday.
Circle the numbers
of any questions
that you are still
unsure how to do.
• I will come around
and grade them at
the end of 5
minutes.
